Electrochemistry

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Electrochemistry

• Lincoln High School

Version 1.03 Updated February 24, 2009
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Electrochemistry

• Electrochemistry is the study of chemical
  reactions that generate electrical effects and of
  the chemical reactions that are caused by the
  action of an electrical current or applied
  potential

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Oxidation and Reduction

• Oxidation loss of electrons.
• Reduction gain of electrons.
• Oxidizing agent substance that causes oxidation
  to occur. The oxidizing agent is reduced.
• Reducing agent substance that causes reduction
  to occur The reducing agent is oxidized.

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Electrical Terminology

• Electrical current may be either direct (DC). or
  alternating (AC).
• In direct current the electrons flow in a single
  direction from negative to positive.
• In an alternating current the direction of
  current flow changes periodically. In the USA
  there are 60 cycles per second. European
  electricity is at 50 cycles.
• A location that has an excess of electrons has a
  negative charge. A location that has a
  deficiency of electrons has a positive charge.
• The rate of current flow is measured in amperes.
• The difference in electrical potential is
  measured in volts.
• The resistance to current flow is measured in
  ohms.

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Electrochemical Cells

• An electrical current can produced from a
  harnessed chemical reaction.
• This system is known as an electrochemical cell
• Voltaic cells are also known as galvanic cells or
  simply as voltaic cells.

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Example 1 The Daniel Cell

• The copper electrode is placed in a solution of
  Cu2 such as copper (II) sulfate or copper (II)
  nitrate.
• The zinc electrode is placed into a solution of
  Zn2 ion such as zinc sulfate or zinc nitrate.
• The two sides are connected with a U tube
  containing an electrolyte such as KCl or KNO3.
  This structure is called a salt bridge.

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Daniel Cell

• Cells and Cell Reactions in a Daniel Cell
• Overall reaction
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• Oxidation half reaction at the anode
• Zn(s) ? Zn2(aq) 2 e-
• Reduction half reaction at the
• cathode
• Cu2(aq) 2 e- ? Cu(s)

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Electrochemical Cells
The voltaic cell on the left has a Potential
difference of about 1.1 volts
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A Standard Hydrogen Electrode

• A hydrogen electrode consists of a platinum
  electrode covered with a fine powder of platinum
  around which H2(g) is bubbled.
• Its potential is defined as zero volts.
• It is the reference point for potential
  measurements
• Hydrogen Half-Cell
• H2(g) ? 2 H(aq) 2 e-
• A reversible reaction

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Distinguishing the Anode and Cathode

• Oxidation occurs at the anode
• Reduction occurs at the cathode

Oxidation (loss of e-) Reduction (gain of e-)
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Distinguishing the Anode and Cathode

• Oxidation occurs at the anode
• Reduction occurs at the cathode

Oxidation (loss of e-) Reduction (gain of e-)
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Standard Reduction Potentials
Zinc when paired with a standard hydrogen
electrode as the cathode produces an electrode
potential of 0.76 volts.
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Standard Reduction Potentials
Copper when paired with a standard hydrogen
electrode as the anode produces an electrode
potential of 0.76 volts.
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Standard Reduction Potentials

• The standard conditions for electrochemical cell
  reactions are
• 25oC
• 1M concentrations for all ions
• 1 atmosphere pressure for all gases
• The standard reduction potential table shows the
  reduction potentials at these conditions relative
  to hydrogen for various reduction of a
  half-reactions

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Standard Reduction Potentials (See Appendix E in
your text, Pages 990-991)

• Li e- ? Li - 3.05 v
• Mg2 2 e- ?Mg - 2.37 v
• Al3 3e- ? Al -1.66 v
• Zn2 2 e- ? Zn - 0.763 v
• Fe2 2 e- ? Fe - 0.440v
• 2 H(aq) 2 e- ? H2(g) 0.00v
• Cu2 2 e- ? Cu 0.337v
• Ag e- ? Ag 0.799v
• O2(g) 4 H(aq) 4 e- ? 2 H2O(l) 1.229v
• F2 2e- ? 2 F- 2.87v

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Application Question 1

• If the reduction of mercury (I) in a voltaic
  cell is desired, the half reaction is
• Which of the following reactions could be used
  as the anode (oxidation)?

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Using Cell Potentials

• They show the potential difference, in volts,
  between the electrodes of an electrochemical
  cell.
• They indicate the direction of Oxidation-Reduction
  reactions.
• A positive value indicates a spontaneous reaction
  indicates that the direction is positive.

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Shorthand Notation for Electrochemical cells

• The shorthand representation of an
  electrochemical cell showing the two half-cells
  connected by a salt bridge or porous barrier,
  such as
• Zn(s)/ZnSO4(aq)//CuSO4(aq)/Cu(s)
• anode cathode

The electrodes are shown on the ends and the
electrolytes for each side are shown in the
middle.
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Calculating Cell Potentials From Standard
Reduction Potenials

• Calculate the cell potential for a cell made from
  silver and zinc electrodes.

From the standard reduction table Zn2 2 e-
? Zn - 0.763 v Ag e- ? Ag 0.799v
Since there must be one oxidation and one
reduction, the direction of one of two half
reactions above must be reversed. Reversing the
zinc half reaction making it the oxidation would
yield a positive cell potential Zn ? Zn2 2
e- 0.763 v (anode) Ag e- ? Ag 0.799v
(cathode) Cell potential 1.562 volts
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Application Question 2

• Mg2 2e- ? Mg E -2.37 V
• Zn2 2e- ? Zn E - 0.76 V
• Does Zn react with Mg2?
• Does Mg react with Zn2?

If Mg is oxidized Mg ? Mg2 2e- Eo 2.37
v Combining this with the reduction of Zn2 Eo
- 0.76 v Leaves an overall positive cell
potential 1.66 v Therefore Mg reacts with
Zn2. Zn does not react with Mg 2
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Metal Displacement Reactions

• The electrochemical cell potentials form the
  basis for predicting which metals will react with
  salt solution of other metals
• This order of reactivity of metals in single
  replacement reactions is called the activity
  series
• The solid of more reactive metals will displace
  ions of a less reactive metal from solution.
• The relative reactivity of metals is based on
  potentials of half reactions.
• Elements with very different potentials react
  most vigorously.

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The Activity Series

• Elements with highly negative reduction
  potentials are not easily reduced but they are
  easily oxidized.
• Since metals react by being oxidized the more
  negative the reduction potential the more
  reactive the element.
• Elements higher in the table (more negative
  potential) can displace any element lower (more
  positive potential).
• So Zn CuCl2 ? ZnCl2 Cu
• Cu ZnCl2 ? No Reaction

The activity Series is really a reduction
potential table arranged from negative to positive
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Gibbs Free Energyand Cell Potential

• DG - nFE RT LnQ
• where n number of electrons changed
• F Faradays constant
• E cell potential

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The Nernst Equation -- Effect of Concentration on
Cell Voltage

• Takes into account corrections for systems that
  are not operating at standard conditions
• Ecell Eocell - (RT/nF) LnQ
• Where R 8.314 J mol-1 K-1
• T Kelvin temperature
• n moles of
  electrons transferred
• F Faradays constant
  96,500 C mol-1
• Q reaction
  quotient
• products/reactants
• and 1 J C-1 1 volt

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The Nernst Equation An Alternate Form

• If the temperature is fixed at 298 K and the
  natural log is replaced with a common log an
  alternate form for the Nernst equation can be
  written as follows
• Ecell Eocell - (0.0591/n)log Q
• Where n moles of electrons transferred
• Q reaction quotient
  
• products/reactants
• The alternate form of the Nernst equation may be
  a little easier to use, but it is less versatile
  since the temperature must be fixed at standard
  thermodynamic temperature

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What is the cell potential for the Daniel's cell
when the Zn2 10 Cu2 ? Assume the
temperature is 25oC. Q (Zn2/Cu2
(10 Cu2)/Cu2 10Eo (0.34 V)Cu couple
(-(-0.76 V)Zn couple 1.10 Voltsand n 2
since 2 electrons are transferred between Zinc
and copper
The Nernst Equation Sample Problem 1

• thus Ecell 1.10 V- (8.314 J mol-1K-1)(298K)
  (Ln10) V
• ( 96,500 C mol-1)(2)
• Ecell 1.100V - 0.0296V
  1.074 V

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Concentration and the Nernst Equation

• In the diagram at the left the half cell
  reactions are the same but the concentrations are
  different
• Will there be electron flow?

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Concentration and the Nernst Equation
Ag Ag e- E1/2 ? V
Anode
Ag e- Ag E1/2 0.80 V
Cathode
Ecell Ecell - (0.0591/n)log(Q)
0 V
1
Ecell - (0.0591) log(0.1) 0.0591 V
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Batteries Are Applications of Electrochemical
Cells

• Batteries
• device that converts chemical energy into
  electricity
• Primary Cells
• non-reversible electrochemical cell
• non-rechargeable cell
• Secondary Cells
• reversible electrochemical cell
• rechargeable cell

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A Common Dry Cell
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A 9 Volt Dry Cell
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Flash Light Batteries

• Dry Cell
• Zn (s) 2 MnO2 (s) 2 NH4 (aq) ?
• Zn2 (aq) 2 MnO(OH) (s) 2 NH3
• Alkaline Cell
• Zn (s) ) 2 MnO2 (s) ? ZnO (s) Mn2O3 (s)

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Lead-Acid (Car Battery)
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Lead-Acid (Car Battery)

• Overall reaction
• Pb (s) PbO2 (s) 2 H2SO4 (aq) 2 PbSO4
  (s) 2 H2O (l)
• E 2.0- volts per cell

Cathode PbO2 (s) SO42- (aq) 4H (aq) 2e- ?
PbSO4 (s) 2 H2O (l) Anode Pb (s) SO42
(aq) ? PbSO4 (s) 2e-
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Nickel-Cadmium (Ni-Cad)

• Overall reaction
• Cd(s) 2 Ni(OH)3(s) Cd(OH)2(s) 2
  Ni(OH)2(s)
• E NiCad 1.25 v/cell

Cathode NiO2 (s) 2 H2O (l) 2e- ? Ni(OH)2 (s)
2OH- (aq) Anode Cd (s) 2OH- (aq) ? Cd(OH)2
(s) 2e-
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Electrolysis
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Electrolysis

• An electrolysis is the inverse of an
  electrochemical cell.
• A non-spontaneous reaction is caused by the
  passage of an electric current through a
  solution.
• By passing a DC current through the an
  electrolyte, the reaction can be made to proceed
  in the reverse or non-spontaneous direction

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Electrolysis

• The reactions at the anode and cathode depend on
  the relative reduction potentials of the solute
  and the solvent.
• The substance produced at the cathode depends on
  the cation that has the higher (more positive)
  reduction potential
• The substance produced at the anode depends on
  the cation that has the lower (more negative)
  reduction potential

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Diagram of a Simple Electrolysis
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Electrolysis of Molten NaCl

• If sodium chloride is heated to its melting
  point, then the resulting liquid contains mobile
  ions. This is a way of producing sodium metal.

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Electrolysis of Brine

• The electrolysis of brine solution results in the
  reduction of water to hydrogen gas rather than
  sodium ion to sodium metal

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Electrolysis of Water

• The electrolysis of water requires a small amount
  of sulfuric acid to be added. Hydrogen and
  oxygen are produced in a 2 to 1 ratio.

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Electrolysis of copper sulfate with a copper
electrodes

• To electroplate a metal, the object to be plated
  is made the cathode and the metal to be plated is
  the anode. The electrolyte is a solution
  containing the cation to be plated.

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Electrolysis Calculations

• The amount of a substance produced during the
  electrolysis reaction depends on the current
  applied and the time the reaction is allowed to
  run
• 1 coulomb 1 ampere second
• 1 mole e- 96,500 coulombs 1 Faraday
• Any combination of current and time that will
  result in 1 Faraday of charge will produce 1 mole
  of electrons.

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Sample Problem 1

• Example 1 How many grams of chromium can be
  plated from a Cr6 solution in 45 minutes at a 25
  amp current?
• (45 min ) (60 s min-1) (25 amp) (1 mol
  e-)(52 g mol-1Cr)
• ----------------------------------------------
  --------------------------
• (96,500 amp s) (6 mol e- mol-1 Cr)

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Sample Problem 1

• Example 1 How many grams of chromium can be
  plated from a Cr6 solution in 45 minutes at a 25
  amp current?
• (45 min ) (60 s min-1) (25 amp) (1 mol e-)(52
  g mol-1Cr)
• ----------------------------------------------
  --------------------------
• (96,500 amp s) (6 mol e- mol-1 Cr)
• 58 g Cr

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Electrolysis Applications

• Preparation of Aluminum (Hall-Heroult process)

• The industrial production of aluminum is
  accomplished by the electrolysis of relatively
  pure alumina
• This process was first invented in France in
  1886 by Paul Heroult and at almost the same time
  in the United States by Charles Hall.
• Adding cryolite, Na3AlF6, to alumina results the
  mixture can be made to melt at 980 C. rather
  than the more than 2000oC of alumina alone.
• It is then electrolyzed using graphite
  electrodes.

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Hall-Heroult Process for Aluminum

• The alumina / cryolite mixture is electrolyzed
  using graphite electrodes.
• Aluminum forms at the cathode and oxygen at the
  anode.
• The oxygen reacts slowly with the carbon anode
  to produce carbon dioxide gas.

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Hall-Heroult Process for Aluminum
Chemical reactions in the processing of aluminium
Alumina reacts with cryolite Al2O3
4 AlF63- ? 3 Al2OF62- 6 F- Cathode
AlF63- 3 e- ? Al 6 F-
Anode 2 Al2OF62- 12 F- C ? 4
AlF63- CO2 4 e-   The overall cell
reaction Al2O3 3 C ? 4 Al
3 CO2
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Hall-Heroult Process for Aluminum
The Hall-Heroult process produces aluminum that
is about 99.5 pure.

• The aluminum produced by the Hall-Heroult process
  is about 99.5 pure. Large quantities of
  electricity are required to produce the aluminum.
  
• Aluminum electrolysis cells operate at a very low
  potential ranging from 4.0 to 5.5 volts but at an
  electrical current of 50,000 to 250,000 amperes.
  
• Each kilogram of aluminum requires between 13 and
  16 kilowatt hours of electrical energy, in
  addition to the energy required to heat the
  alumina/cryolite mixture.  

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Chlor-Alkali Processes

• Electrolysis of Sodium chloride --
• With molten sodium chloride the products are
  liquid sodium and chlorine gas
• With aqueous sodium chloride or brine the
  products are sodium hydroxide (caustic soda) and
  chlorine gas.

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Electrolysis of Molten NaCl
The electrolysis of molten NaCl at high
temperatures generates liquid sodium metal and
chlorine gas.
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Industrial Electrolysis of Brine
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Corrosion

• Corrosion of metals is a common
  oxidation-reduction process in nature.
• The rusting of iron can be thought of as a form
  of an electrochemical cell.

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Rusting of Iron
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Rusting Iron

• O2(g) 4 H(aq) 4 e- ? 2 H2O(l)
  Eo 1.23V
• Rusting Process
• Fe(s) ? Fe2(aq) 2 e- Eo
  0.44 V
• O2(g) 4 H(aq) 4 e- ? 2 H2O(l) Eo
  1.23 V
• --------------------------------------------------
  ------- --------------
• 2 Fe(s) O2(g) 4 H(aq) ?2 H2O(l) Fe2(aq)
  Eo 1.67 V

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Preventing Corrosion

• painting
• galvanizing
• sacrificial anode

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